Topic 3 - Bonding: Valence Bond Theory (VBT)

Chemical Bond

A short range interaction between two or more species which result in the formation of a stable molecular identity.

Valence Bond Theory (VBT)

• Core assumption: Electrons occupy atomic orbitals; bonds form via orbital overlap between two atoms (electron donation from each atom).

• Orbital overlap as waves: Orbital overlap means that the two waves (atoms are in-phase → in-phase overlap increases amplitude → high electron density between nuclei.

Sigma (σ) Bonds

• Formation: Head-on overlap of atomic orbitals (s–s, s–p, or p–p).

• Shape: Cylindrically symmetrical around the bond axis that passes through both nuclei.

• Rule: Only one sigma bond can exist between any two atoms.

  • The second that there's like-charges that repel electrons away from each other, further orbital overlap (bonding cannot occur in the same location.

• Stronger Bond:

  • Type of Overlap: A head-to-head overlap maximizes the electron density.

Pi (π) Bonds

• Formation: Side-to-side overlap of p orbitals (above and below bond axis).

• Characteristics:

  • Two regions of electron density (above and below nuclei).

  • Still considered one bond despite two regions.

• Weaker Bond:

  • Type of Overlap:

    • A side-to-side overlap means less orbital overlap which minimizes the electron density.

    • Since electron densities are spread above and below the bond axis, there's weaker electrostatic attraction between the nuclei and bonding electrons

  • Orbital Alignment: A side-to-side overlap is more restricted in orientation. Any twisting/rotation would misalign the orbitals.

Single Bond

1 σ

Double Bond

1 σ + 1 π

Triple bond

1 σ + 2 π

Atomic Number

• Number of protons (+) in the nucleus (atomic number) determine what element it is.

• The nucleus of an atom has an overall positive charge equal to the number of proton it contains.

Electrostatic Interactions

When two atoms come together to form a chemical bond, many electrostatic interactions happen simulatenously:

• Nucleus (+) ↔ electrons (-) → attraction.

• Nucleus (+) ↔ nucleus (+) → repulsion.

• Electron (-) ↔ electron (-) → repulsion.

Electrostatic Interactions - Process

 

1. No interaction between atoms, the inter-nuclear distance is too large.

2. Atoms start coming closer together → nuclei attracted to each other's electrons → attractive forces dominate and outweighs repulsive forces → potential energy keeps dropping until a minimum.

3. Bonding occurs at a distance where the potential energy reaches a minimum point - point of maximum stability.

4. Once atoms get too close together → nucleus–nucleus repulsion dominates → energy rises sharply.

*Zero distance is impossible, nuclear repulsion is infinitely strong!

Total Potential Energy

Energy of the interaction of two atoms; all attractive forces + all repulsive forces

Bond Length

Internuclear distance at minimum energy - where bonding occurs.

Bond Energy

Energy needed to break a chemical bond, that is the energy difference between non-interacting atoms and atoms at minimum energy; always positive.

Bond Order

• Definition: Number of chemical bonds between a pair of atoms.

• Relationships: Higher bond order → shorter bond length → higher bond energy → greater stability.

• Trends: Smaller atoms (e.g., H₂) → shorter bonds, as the valence electrons are in orbitals that are physically closer to the nucleus.